The Structure of Atoms

1
CHAPTER 6

• The Structure of Atoms

2
Chapter Outline

• Subatomic Particles
• Fundamental Particles
• The Discovery of Electrons
• Canal Rays and Protons
• Rutherford and the Nuclear Atom
• Atomic Number
• Neutrons
• Mass Number and Isotopes
• Mass spectrometry and Isotopic Abundance

3
Chapter Goals

• The Atomic Weight Scale and Atomic Weights
• The Electronic Structures of Atoms
• Electromagnetic radiation
• The Photoelectric Effect
• Atomic Spectra and the Bohr Atom
• The Wave Nature of the Electron
• The Quantum Mechanical Picture of the Atom

4
Chapter Goals

• Quantum Numbers
• Atomic Orbitals
• Electron Configurations
• Paramagnetism and Diamagnetism
• The Periodic Table and Electron Configurations

5
Fundamental Particles

• Three fundamental particles make up atoms. The
  following table lists these particles together
  with their masses and their charges.

6
The Discovery of Electrons

• Humphrey Davy in the early 1800s passed
  electricity through compounds and noted
• that the compounds decomposed into elements.
• Concluded that compounds are held together by
  electrical forces.
• Michael Faraday in 1832-1833 realized that the
  amount of reaction that occurs during
  electrolysis is proportional to the electrical
  current passed through the compounds.

7
The Discovery of Electrons

• Cathode Ray Tubes experiments performed in the
  late 1800s early 1900s.
• Consist of two electrodes sealed in a glass tube
  containing a gas at very low pressure.
• When a voltage is applied to the cathodes a glow
  discharge is emitted.

8
The Discovery of Electrons

• These rays are emitted from cathode (- end) and
  travel to anode ( end).
• Cathode Rays must be negatively charged!
• J.J. Thomson modified the cathode ray tube
  experiments in 1897 by adding two adjustable
  voltage electrodes.
• Studied the amount that the cathode ray beam was
  deflected by additional electric field.

9
The Discovery of Electrons

• Modifications to the basic cathode ray tube
  experiment.

10
The Discovery of Electrons

• Thomson used his modification to measure the
  charge to mass ratio of electrons.
• Charge to mass ratio
• e/m -1.75881 x 108 coulomb/g of e-
• Thomson named the cathode rays electrons.
• Thomson is considered to be the discoverer of
  electrons.
• TV sets and computer screens are cathode ray
  tubes.

11
The Discovery of Electrons

• Robert A. Millikan won the 1st American Nobel
  Prize in 1923 for his famous oil-drop experiment.
• In 1909 Millikan determined the charge and mass
  of the electron.

12
The Discovery of Electrons

• Millikan determined that the charge on a single
  electron -1.60218 x 10-19 coulomb.
• Using Thomsons charge to mass ratio we get that
  the mass of one electron is 9.11 x 10-28 g.
• e/m -1.75881 x 108 coulomb
• e -1.60218 x 10-19 coulomb
• Thus m 9.10940 x 10-28 g

13
Canal Rays and Protons

• Eugene Goldstein noted streams of positively
  charged particles in cathode rays in 1886.
• Particles move in opposite direction of cathode
  rays.
• Called Canal Rays because they passed through
  holes (channels or canals) drilled through the
  negative electrode.
• Canal rays must be positive.
• Goldstein postulated the existence of a positive
  fundamental particle called the proton.

14
Rutherford and the Nuclear Atom

• Ernest Rutherford directed Hans Geiger and Ernst
  Marsdens experiment in 1910.
• ?- particle scattering from thin Au foils
• Gave us the basic picture of the atoms
  structure.

15
Rutherford and the Nuclear Atom

• In 1912 Rutherford decoded the ?-particle
  scattering information.
• Explanation involved a nuclear atom with
  electrons surrounding the nucleus .

16
Rutherford and the Nuclear Atom

• Rutherfords major conclusions from the
  ?-particle scattering experiment
• The atom is mostly empty space.
• It contains a very small, dense center called the
  nucleus.
• Nearly all of the atoms mass is in the nucleus.
• The nuclear diameter is 1/10,000 to 1/100,000
  times less than atoms radius.

17
Rutherford and the Nuclear Atom

• Because the atoms mass is contained in such a
  small volume
• The nuclear density is 1015g/mL.
• This is equivalent to 3.72 x 109 tons/in3.
• Density inside the nucleus is almost the same as
  a neutron stars density.

18
Atomic Number

• The atomic number is equal to the number of
  protons in the nucleus.
• Sometimes given the symbol Z.
• On the periodic chart Z is the uppermost number
  in each elements box.
• In 1913 H.G.J. Moseley realized that the atomic
  number determines the element .
• The elements differ from each other by the number
  of protons in the nucleus.
• The number of electrons in a neutral atom is also
  equal to the atomic number.

19
Neutrons

• James Chadwick in 1932 analyzed the results of
  ?-particle scattering on thin Be films.
• Chadwick recognized existence of massive neutral
  particles which he called neutrons.
• Chadwick discovered the neutron.

20
Mass Number and Isotopes

• Mass number is given the symbol A.
• A is the sum of the number of protons and
  neutrons.
• Z proton number N neutron number
• A Z N
• A common symbolism used to show mass and proton
  numbers is

• Can be shortened to this symbolism.

21
Mass Number and Isotopes

• Isotopes are atoms of the same element but with
  different neutron numbers.
• Isotopes have different masses and A values but
  are the same element.
• One example of an isotopic series is the hydrogen
  isotopes.
• 1H or protium is the most common hydrogen
  isotope.
• one proton and no neutrons
• 2H or deuterium is the second most abundant
  hydrogen isotope.
• one proton and one neutron
• 3H or tritium is a radioactive hydrogen isotope.
• one proton and two neutrons

22
Mass Number and Isotopes

• The stable oxygen isotopes provide another
  example.
• 16O is the most abundant stable O isotope.
• How many protons and neutrons are in 16O?

• 17O is the least abundant stable O isotope.
• How many protons and neutrons are in 17O?

• 18O is the second most abundant stable O
  isotope.
• How many protons and neutrons in 18O?

23
Mass Spectrometry andIsotopic Abundances

• Francis Aston devised the first mass
  spectrometer.
• Device generates ions that pass down an evacuated
  path inside a magnet.
• Ions are separated based on their mass.

24
Mass Spectrometry andIsotopic Abundances

• There are four factors which determine a
  particles path in the mass spectrometer.
• accelerating voltage
• magnetic field strength
• masses of particles
• charge on particles

25
Mass Spectrometry andIsotopic Abundances

• Mass spectrum of Ne ions shown below.
• How scientists determine the masses and
  abundances of the isotopes of an element.

26
The Atomic Weight Scale and Atomic Weights

• If we define the mass of 12C as exactly 12 atomic
  mass units (amu), then it is possible to
  establish a relative weight scale for atoms.
• 1 amu (1/12) mass of 12C by definition
• What is the mass of an amu in grams?
• Example 5-1 Calculate the number of atomic mass
  units in one gram.
• The mass of one 31P atom has been experimentally
  determined to be 30.99376 amu.
• 1 mol of 31P atoms has a mass of 30.99376 g.

27
The Atomic Weight Scale and Atomic Weights
28
The Atomic Weight Scale and Atomic Weights

• Thus 1.00 g 6.022 x 1023 amu.
• This is always true and provides the conversion
  factor between grams and amu.

29
The Atomic Weight Scale and Atomic Weights

• The atomic weight of an element is the weighted
  average of the masses of its stable isotopes
• Example 5-2 Naturally occurring Cu consists of 2
  isotopes. It is 69.1 63Cu with a mass of 62.9
  amu, and 30.9 65Cu, which has a mass of 64.9
  amu. Calculate the atomic weight of Cu to one
  decimal place.

30
The Atomic Weight Scale and Atomic Weights
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The Atomic Weight Scale and Atomic Weights
32
The Atomic Weight Scale and Atomic Weights
33
The Atomic Weight Scale and Atomic Weights

• Example 5-3 Naturally occurring chromium
  consists of four isotopes. It is 4.31 2450Cr,
  mass 49.946 amu, 83.76 2452Cr, mass 51.941
  amu, 9.55 2453Cr, mass 52.941 amu, and 2.38
  2454Cr, mass 53.939 amu. Calculate the atomic
  weight of chromium.
• You do it!

34
The Atomic Weight Scale and Atomic Weights
35
The Atomic Weight Scale and Atomic Weights

• Example 5-4 The atomic weight of boron is 10.811
  amu. The masses of the two naturally occurring
  isotopes 510B and 511B, are 10.013 and 11.009
  amu, respectively. Calculate the fraction and
  percentage of each isotope.
• You do it!
• This problem requires a little algebra.
• A hint for this problem is x (1-x) 1

36
The Atomic Weight Scale and Atomic Weights
37
The Atomic Weight Scale and Atomic Weights

• Note that because x is the multiplier for the 10B
  isotope, our solution gives us the fraction of
  natural B that is 10B.
• Fraction of 10B 0.199 and abundance of 10B
  19.9.
• The multiplier for 11B is (1-x) thus the fraction
  of 11B is 1-0.199 0.811 and the abundance of
  11B is 81.1.

38
The Electronic Structures of AtomsElectromagnetic
Radiation

• The wavelength of electromagnetic radiation has
  the symbol ??.
• Wavelength is the distance from the top (crest)
  of one wave to the top of the next wave.
• Measured in units of distance such as m,cm, Å.
• 1 Å 1 x 10-10 m 1 x 10-8 cm
• The frequency of electromagnetic radiation has
  the symbol ?.
• Frequency is the number of crests or troughs that
  pass a given point per second.
• Measured in units of 1/time - s-1

39
Electromagnetic Radiation

• The relationship between wavelength and frequency
  for any wave is velocity ???.
• For electromagnetic radiation the velocity is
  3.00 x 108 m/s and has the symbol c.
• Thus c ??? for electromagnetic radiation.

40
Electromagnetic Radiation

• Molecules interact with electromagnetic
  radiation.
• Molecules can absorb and emit light.
• Once a molecule has absorbed light (energy), the
  molecule can
• Rotate
• Translate
• Vibrate
• Electronic transition

41
Electromagnetic Radiation

• For water
• Rotations occur in the microwave portion of
  spectrum.
• Vibrations occur in the infrared portion of
  spectrum.
• Translation occurs across the spectrum.
• Electronic transitions occur in the ultraviolet
  portion of spectrum.

42
Electromagnetic Radiation

• Example 5-5 What is the frequency of green light
  of wavelength 5200 Å?

43
Electromagnetic Radiation

• In 1900 Max Planck studied black body radiation
  and realized that to explain the energy spectrum
  he had to assume that
• energy is quantized
• light has particle character
• Plancks equation is

44
Electromagnetic Radiation

• Example 5-6 What is the energy of a photon of
  green light with wavelength 5200 Å? What is the
  energy of 1.00 mol of these photons?

45
The Photoelectric Effect

• Light can strike the surface of some metals
  causing an electron to be ejected.

46
The Photoelectric Effect

• What are some practical uses of the photoelectric
  effect?
• You do it!
• Electronic door openers
• Light switches for street lights
• Exposure meters for cameras
• Albert Einstein explained the photoelectric
  effect
• Explanation involved light having particle-like
  behavior.
• Einstein won the 1921 Nobel Prize in Physics for
  this work.

47
Atomic Spectra and the Bohr Atom

• An emission spectrum is formed by an electric
  current passing through a gas in a vacuum tube
  (at very low pressure) which causes the gas to
  emit light.
• Sometimes called a bright line spectrum.

48
Atomic Spectra and the Bohr Atom

• An absorption spectrum is formed by shining a
  beam of white light through a sample of gas.
• Absorption spectra indicate the wavelengths of
  light that have been absorbed.

49
Atomic Spectra and the Bohr Atom

• Every element has a unique spectrum.
• Thus we can use spectra to identify elements.
• This can be done in the lab, stars, fireworks,
  etc.

50
Atomic Spectra and the Bohr Atom

• Atomic and molecular spectra are important
  indicators of the underlying structure of the
  species.
• In the early 20th century several eminent
  scientists began to understand this underlying
  structure.
• Included in this list are
• Niels Bohr
• Erwin Schrodinger
• Werner Heisenberg

51
Atomic Spectra and the Bohr Atom

• Example 5-7 An orange line of wavelength 5890 Å
  is observed in the emission spectrum of sodium.
  What is the energy of one photon of this orange
  light?
• You do it!

52
Atomic Spectra and the Bohr Atom

• The Rydberg equation is an empirical equation
  that relates the wavelengths of the lines in the
  hydrogen spectrum.

53
Atomic Spectra and the Bohr Atom

• Example 5-8. What is the wavelength of light
  emitted when the hydrogen atoms energy changes
  from n 4 to n 2?

54
Atomic Spectra and the Bohr Atom
Notice that the wavelength calculated from the
Rydberg equation matches the wavelength of the
green colored line in the H spectrum.
55
Atomic Spectra and the Bohr Atom

• In 1913 Neils Bohr incorporated Plancks quantum
  theory into the hydrogen spectrum explanation.
• Here are the postulates of Bohrs theory.

• Atom has a number of definite and discrete energy
  levels (orbits) in which an electron may exist
  without emitting or absorbing electromagnetic
  radiation.
• As the orbital radius increases so does the
  energy
• 1lt2lt3lt4lt5......

56
Atomic Spectra and the Bohr Atom

• An electron may move from one discrete energy
  level (orbit) to another, but, in so doing,
  monochromatic radiation is emitted or absorbed in
  accordance with the following equation.

• Energy is absorbed when electrons jump to higher
  orbits.
• n 2 to n 4 for example
• Energy is emitted when electrons fall to lower
  orbits.
• n 4 to n 1 for example

57
Atomic Spectra and the Bohr Atom

• An electron moves in a circular orbit about the
  nucleus and it motion is governed by the ordinary
  laws of mechanics and electrostatics, with the
  restriction that the angular momentum of the
  electron is quantized (can only have certain
  discrete values).
• angular momentum mvr nh/2?
• h Plancks constant n 1,2,3,4,...(energy
  levels)
• v velocity of electron m mass of electron
• r radius of orbit

58
Atomic Spectra and the Bohr Atom

• Light of a characteristic wavelength (and
  frequency) is emitted when electrons move from
  higher E (orbit, n 4) to lower E (orbit, n
  1).
• This is the origin of emission spectra.

• Light of a characteristic wavelength (and
  frequency) is absorbed when electron jumps from
  lower E (orbit, n 2) to higher E (orbit, n 4)
• This is the origin of absorption spectra.

59
Atomic Spectra and the Bohr Atom

• Bohrs theory correctly explains the H emission
  spectrum.
• The theory fails for all other elements because
  it is not an adequate theory.

60
The Wave Nature of the Electron

• In 1925 Louis de Broglie published his Ph.D.
  dissertation.
• A crucial element of his dissertation is that
  electrons have wave-like properties.
• The electron wavelengths are described by the de
  Broglie relationship.

61
The Wave Nature of the Electron

• De Broglies assertion was verified by Davisson
  Germer within two years.
• Consequently, we now know that electrons (in fact
  - all particles) have both a particle and a wave
  like character.
• This wave-particle duality is a fundamental
  property of submicroscopic particles.

62
The Wave Nature of the Electron

• Example 5-9. Determine the wavelength, in m, of
  an electron, with mass 9.11 x 10-31 kg, having a
  velocity of 5.65 x 107 m/s.
• Remember Plancks constant is 6.626 x 10-34 Js
  which is also equal to 6.626 x 10-34 kg m2/s2.

63
The Wave Nature of the Electron

• Example 5-10. Determine the wavelength, in m, of
  a 0.22 caliber bullet, with mass 3.89 x 10-3 kg,
  having a velocity of 395 m/s, 1300 ft/s.
• You do it!

• Why is the bullets wavelength so small compared
  to the electrons wavelength?

64
The Quantum Mechanical Picture of the Atom

• Werner Heisenberg in 1927 developed the concept
  of the Uncertainty Principle.
• It is impossible to determine simultaneously both
  the position and momentum of an electron (or any
  other small particle).
• Detecting an electron requires the use of
  electromagnetic radiation which displaces the
  electron!
• Electron microscopes use this phenomenon

65
The Quantum Mechanical Picture of the Atom

• Consequently, we must speak of the electrons
  position about the atom in terms of probability
  functions.
• These probability functions are represented as
  orbitals in quantum mechanics.

66
The Quantum Mechanical Picture of the Atom

• Basic Postulates of Quantum Theory
• Atoms and molecules can exist only in certain
  energy states. In each energy state, the atom or
  molecule has a definite energy. When an atom or
  molecule changes its energy state, it must emit
  or absorb just enough energy to bring it to the
  new energy state (the quantum condition).

67
The Quantum Mechanical Picture of the Atom

• Atoms or molecules emit or absorb radiation
  (light) as they change their energies. The
  frequency of the light emitted or absorbed is
  related to the energy change by a simple
  equation.

68
The Quantum Mechanical Picture of the Atom

• The allowed energy states of atoms and molecules
  can be described by sets of numbers called
  quantum numbers.
• Quantum numbers are the solutions of the
  Schrodinger, Heisenberg Dirac equations.
• Four quantum numbers are necessary to describe
  energy states of electrons in atoms.

69
Quantum Numbers

• The principal quantum number has the symbol n.
• n 1, 2, 3, 4, ...... shells
• n K, L, M, N, ......
• The electrons energy depends principally on n .

70
Quantum Numbers

• The angular momentum quantum number has the
  symbol ?.
• ? 0, 1, 2, 3, 4, 5, .......(n-1)
• ? s, p, d, f, g, h, .......(n-1)
• ? tells us the shape of the orbitals.
• These orbitals are the volume around the atom
  that the electrons occupy 90-95 of the time.
• This is one of the places where Heisenbergs
  Uncertainty principle comes into play.

71
Quantum Numbers

• The symbol for the magnetic quantum number is m?.
• m? - ? , (- ? 1), (- ? 2), .....0, .......,
  (? -2), (? -1), ?
• If ? 0 (or an s orbital), then m? 0.
• Notice that there is only 1 value of m?.
• This implies that there is one s orbital per n
  value. n ? 1
• If ? 1 (or a p orbital), then m? -1,0,1.
• There are 3 values of m?.
• Thus there are three p orbitals per n value. n ? 2

72
Quantum Numbers

• If ? 2 (or a d orbital), then m?
  -2,-1,0,1,2.
• There are 5 values of m?.
• Thus there are five d orbitals per n value. n ?
  3
• If ? 3 (or an f orbital), then m?
  -3,-2,-1,0,1,2, 3.
• There are 7 values of m?.
• Thus there are seven f orbitals per n value, n
• Theoretically, this series continues on to g,h,i,
  etc. orbitals.
• Practically speaking atoms that have been
  discovered or made up to this point in time only
  have electrons in s, p, d, or f orbitals in their
  ground state configurations.

73
Quantum Numbers

• The last quantum number is the spin quantum
  number which has the symbol ms.
• The spin quantum number only has two possible
  values.
• ms 1/2 or -1/2
• ms 1/2
• This quantum number tells us the spin and
  orientation of the magnetic field of the
  electrons.
• Wolfgang Pauli in 1925 discovered the Exclusion
  Principle.
• No two electrons in an atom can have the same set
  of 4 quantum numbers.

74
Atomic Orbitals

• Atomic orbitals are regions of space where the
  probability of finding an electron about an atom
  is highest.
• s orbital properties
• There is one s orbital per n level.
• ? 0 1 value of m?

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Atomic Orbitals

• s orbitals are spherically symmetric.

76
Atomic Orbitals

• p orbital properties
• The first p orbitals appear in the n 2 shell.
• p orbitals are peanut or dumbbell shaped volumes.
• They are directed along the axes of a Cartesian
  coordinate system.
• There are 3 p orbitals per n level.
• The three orbitals are named px, py, pz.
• They have an ? 1.
• m? -1,0,1 3 values of m?

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Atomic Orbitals

• p orbitals are peanut or dumbbell shaped.

78
Atomic Orbitals

• d orbital properties
• The first d orbitals appear in the n 3 shell.
• The five d orbitals have two different shapes
• 4 are clover leaf shaped.
• 1 is peanut shaped with a doughnut around it.
• The orbitals lie directly on the Cartesian axes
  or are rotated 45o from the axes.

• There are 5 d orbitals per n level.
• The five orbitals are named
• They have an ? 2.
• m? -2,-1,0,1,2 5 values of m ?

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Atomic Orbitals

• d orbital shapes

80
Atomic Orbitals

• f orbital properties
• The first f orbitals appear in the n 4 shell.
• The f orbitals have the most complex shapes.
• There are seven f orbitals per n level.
• The f orbitals have complicated names.
• They have an ? 3
• m? -3,-2,-1,0,1,2, 3 7 values of m?
• The f orbitals have important effects in the
  lanthanide and actinide elements.

81
Atomic Orbitals

• f orbital shapes

82
Atomic Orbitals

• Spin quantum number effects
• Every orbital can hold up to two electrons.
• Consequence of the Pauli Exclusion Principle.
• The two electrons are designated as having
• one spin up ? and one spin down??
• Spin describes the direction of the electrons
  magnetic fields.

83
Paramagnetism and Diamagnetism

• Unpaired electrons have their spins aligned ?? or
  ??
• This increases the magnetic field of the atom.
• Atoms with unpaired electrons are called
  paramagnetic .
• Paramagnetic atoms are attracted to a magnet.

84
Paramagnetism and Diamagnetism

• Paired electrons have their spins unaligned ??.
• Paired electrons have no net magnetic field.
• Atoms with paired electrons are called
  diamagnetic.
• Diamagnetic atoms are repelled by a magnet.

85
Paramagnetism and Diamagnetism

• Because two electrons in the same orbital must be
  paired, it is possible to calculate the number of
  orbitals and the number of electrons in each n
  shell.
• The number of orbitals per n level is given by
  n2.
• The maximum number of electrons per n level is
  2n2.
• The value is 2n2 because of the two paired
  electrons.

86
Paramagnetism and Diamagnetism

• Energy Level of Orbitals Max. of e-
• n n2 2n2
• 1 1 2
• 2 4 8
• You do it!

3 9 18 4 16 32
87
The Periodic Table and Electron Configurations

• The principle that describes how the periodic
  chart is a function of electronic configurations
  is the Aufbau Principle.
• The electron that distinguishes an element from
  the previous element enters the lowest energy
  atomic orbital available.

88
The Periodic Table and Electron Configurations

• The Aufbau Principle describes the electron
  filling order in atoms.

89
The Periodic Table and Electron Configurations

• There are two ways to remember the correct
  filling order for electrons in atoms.
• You can use this mnemonic.

90
The Periodic Table and Electron Configurations

1. Or you can use the periodic chart .

91
The Periodic Table and Electron Configurations

• Now we will use the Aufbau Principle to determine
  the electronic configurations of the elements on
  the periodic chart.
• 1st row elements.

92
The Periodic Table and Electron Configurations

• 2nd row elements.

• Hunds rule tells us that the electrons will fill
  the
• p orbitals by placing electrons in each orbital
• singly and with same spin until half-filled.
  Then
• the electrons will pair to finish the p orbitals.

93
The Periodic Table and Electron Configurations

• 3rd row elements

94
The Periodic Table and Electron Configurations

• 4th row elements

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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
111
The Periodic Table and Electron Configurations
112
The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
118
The Periodic Table and Electron Configurations

• Now we can write a complete set of quantum
  numbers for all of the electrons in these three
  elements as examples.
• Na
• Ca
• Fe
• First for 11Na.
• When completed there must be one set of 4 quantum
  numbers for each of the 11 electrons in Na
• (remember Ne has 10 electrons)

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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations

• Next we will do the same exercise for 20Ca.
• Again, when finished we must have one set of 4
  quantum numbers for each of the 20 electrons in
  Ca.
• We represent the first 18 electrons in Ca with
  the symbol Ar.

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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations

• Finally, we do the same exercise for 26Fe.
• We should have one set of 4 quantum numbers for
  each of the 26 electrons in Fe.
• To save time and space, we use the symbol Ar to
  represent the first 18 electrons in Fe

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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
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The Periodic Table and Electron Configurations
144
Synthesis Question

• What is the atomic number of the element that
  should theoretically be the noble gas below Rn?
• The 6 ds are completed with element 112 and the
  7ps are completed with element 118. Thus the
  next noble gas (or perhaps it will be a noble
  liquid) should be element 118.

145
Group Question

• In a universe different from ours, the laws of
  quantum mechanics are the same as ours with one
  small change. Electrons in this universe have
  three spin states, -1, 0, and 1, rather than the
  two, 1/2 and -1/2, that we have. What two
  elements in this universe would be the first and
  second noble gases? (Assume that the elements in
  this different universe have the same symbols as
  in ours.)

146
End of Chapter 5

• The study of various spectra is one of the
  fundamental tools that chemists apply to numerous
  areas of their work.